Group 1A — The Alkali Metals

1A 2A 3A 4A 5A 6A 7A 8A
(1) (2) (13) (14) (15) (16) (17) (18)
3B 4B 5B 6B 7B 8B 1B 2B
(3) (4) (5) (6) (7) (8) (9) (10) (11) (12)
1 H He
2 Li Be B C N O F Ne
3 Na Mg Al Si P S Cl Ar
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6 Cs Ba La   Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
7 Fr Ra Ac   Rf Db Sg Bh Hs Mt Ds Rg Uub Uuq
6   Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
7   Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr


Group 1A (or IA) of the periodic table are the alkali metals:  hydrogen (H), lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr).  These are (except for hydrogen) soft, shiny, low-melting, highly reactive metals, which tarnish when exposed to air.  The name comes from the fact that when these metals or their oxides are dissolved in water, a basic (alkaline) solution results.  Because the alkali metals are very reactive, they are seldom (if ever) found in their elemental form in nature, and are usually found as ionic compounds (except for hydrogen).

The alkali metals have only one valence electron in their highest-energy orbitals (ns1).  In their respective periods, they are the largest elements and have the lowest ionization energies.  The valence electron is easily lost, forming an ion with a 1+ charge.

The alkali metals are solids at room temperature (except for hydrogen), but have fairly low melting points:  lithium melts at 181C, sodium at 98C, potassium at 63C, rubidium at 39C, and cesium at 28C.  They are also relatively soft metals:  sodium and potassium can be cut with a butter knife.

Salts of the Group 1A elements tend to be extremely soluble in water.  Because the alkali metal ions are relatively large (compared to other ions from the same period), their charges densities are low, and they are easily separated from their anions and solvated by polar solvents like water.

The alkali metals (again, except for hydrogen) react vigorously with water, producing the metal hydroxide, hydrogen gas, and heat.

2M(s)  +  H2O(l)    MOH(aq)  +  H2(g)

(Heat plus hydrogen in an oxygen atmosphere is, of course, a very dangerous combination!)  The reaction becomes more vigorous as one moves from top to bottom in Group 1A:  lithium sizzles fiercely in water, a small amount of sodium reacts even more vigorously, and even a small amount of potassium metal reacts violently and usually ignites the hydrogen gas; rubidium and cesium explode.  This is a result of the fact that the size of the element increases as we move down the group:  as the size of the metal increases, the valence electron is farther away from the nucleus, and is thus more easily removed (i.e., the ionization energy is lower).


Hydrogen (H, Z=1).

Although hydrogen is placed at the top of Group 1A in most versions of the periodic table, it is very different from the other members of the alkali metal group.  In its elemental form, hydrogen is a colorless, odorless, extremely flammable gas at room temperature, consisting of diatomic molecules of H2.  Molecular hydrogen boils at -253C (20 K), and freezes at -259C (14 K).  Under tremendous pressure (about 2 million atmospheres), it can be converted to a metallic form, capable of conducting electricity.  (It has been theorized that center of the planet Jupiter consists of metallic hydrogen.)  In the Earth's crust, it is found at a concentration of 1500 ppm (mostly in the form water and of organic compounds), making it the 10th most abundant element.

Hydrogen is the most abundant element in the universe (75% by weight, or 88% of all of the atoms of the universe); hydrogen and helium together make up 99% of the "normal" matter of the universe.  (Of course, there's also "dark matter" and "dark energy" to worry about, but that's another story.)  Hydrogen, helium, and trace amounts of lithium were produced at the beginning of the Universe in the Big Bang, and became concentrated into stars by the force of gravity.  The fusion of hydrogen to form helium provides the power that makes stars shine:  in the Sun, 600 millions tons of hydrogen undergo fusion to form helium every second, converting 5 million tons of matter into energy (Einstein's good ol' E = mc2).  The fusion of hydrogen and its isotopes (see below) also powers the hydrogen bomb, which contains lithium deuteride (LiD) and tritium; the explosion of a fission-powered bomb produces neutrons which initiate fusion of the deuterium with the tritium, releasing vast amounts of energy.  Research into achieving controlled nuclear fusion to generate electricity is being conducted, but the extremely high temperatures that are necessary to initiate the fusion reactions present a major challenge to physicists.

Hydrogen typically does not form cations, but instead forms compounds through covalent bonding.  Hydrogen can form bonds to many other elements, such as nitrogen (NH3 and its derivatives), oxygen (H2O) and sulfur (H2S), the halogens (HX), and carbon, where it is found in millions of different hydrocarbons and other organic molecules (almost all organic molecules contain at least some hydrogen atoms).  Hydrogen can also bond to metal atoms, such as lithium (LiH), calcium (CaH2), etc.  In these compounds, the bonding is usually pictured as a metal cation combined with a hydride anion (H-).  (On some periodic tables, in fact, hydrogen is placed at the top of Group 7A, since like the halogens, it can form a -1 charge.)  Hydrogen is also found in acids, which are molecules containing easily-removed hydrogen atoms, usually connected to oxygen, nitrogen, or a halogen.  When dissolved in water, these substances transfer hydrogen as "H+" (often referred to as a proton) to water, forming the hydronium ion, H3O+.  (This is a greatly oversimplified explanation of acid-base chemistry.)  Some commonly encountered acids include hydrochloric acid (HCl, also known a muriatic acid), sulfuric acid (H2SO4), nitric acid (HNO3), acetic acid (HC2H3O2, the active component of vinegar), phosphoric acid (H3PO4), hydrofluoric acid (HF), and many others.

Hydrogen was discovered by the English chemist Henry Cavendish in 1766; hydrogen had been observed before, but Cavendish was the first to recognize not only that it was an element, but that it burned to form water, which also provided conclusive proof that water was not itself an element.  The name "hydrogen" was derived by the French chemist Antoine Lavoisier from the Greek words hydro ("water") and genes ("forming")

There are three isotopes of hydrogen.  Hydrogen-1, or protium, contains one proton in its nucleus, and is by far the most common form of hydrogen (99.985% of all the world's hydrogen).  Hydrogen-2, or deuterium, contains one proton and one neutron in its nucleus, and comprises the remaining 0.015% of the world's naturally-occurring hydrogen.  Hydrogen-3, or tritium, contains one proton and two neutrons, and is only found in trace amounts; it is produced by the interaction of cosmic rays on gases in the upper atmosphere, and in nuclear explosions, but since it has a half life of only 12.3 years, it does not accumulate in the atmosphere.

Heavy water is water made from two atoms of deuterium and one atom of oxygen.  This form of water is literally heavier than "ordinary" water, since an atom of deuterium is twice as heavy as an atom of "regular" hydrogen.  (H2O has a molar mass of 18.02 g/mol; D2O has a molar mass of 20.03 g/mol.)  Ordinary water contains about 1 molecule of D2O for every 7000 molecules of H2O.  The electrolysis of water concentrates D2O in the solution, since the lighter isotope evaporates from the solution slightly faster.  Successive electrolysis experiments allow pure heavy water to be produced, but it takes about 100,000 gallons of water to produce 1 gallon of heavy water by this method.  Heavy water is used as a moderator in nuclear reactions:  it slows down fast-moving neutrons, allowing them to be captured more easily by other nuclei.  The generation of heavy water was important during the research on nuclear fission that went into the Manhattan Project during World War II.  Because the deuterium in heavy water is heavier than ordinary hydrogen, the consumption of heavy water disrupts some cellular processes, especially those that rely heavily on hydrogen bonding (see below):  seeds grown in heavy water do not germinate, and rats die after a week of drinking nothing but heavy water, when their body water reaches 50% deuteration.  (For a typical person, a fatal dose would require drinking nothing but heavy water for 10 to 14 days, so it's pretty doubtful that heavy water poisoning will be featured on CSI anytime soon.)

Most hydrogen is prepared industrially be reacting coal or hydrocarbons with steam at high temperatures to produce carbon monoxide and hydrogen gas (a mixture of carbon monoxide and hydrogen is called synthesis gas, and can be used in manufacturing methanol).  On smaller scales it can be produced by the reaction of active metals (such as zinc, calcium, etc.) with hydrochloric acid, or by the electrolysis of water.

Hydrogen gas is combined with nitrogen in the Haber process to synthesize ammonia (NH3), which is widely used in fertilizers.  It is also used in the manufacture of hydrogenated vegetable oils; in this reaction, hydrogen atoms add to the carbon-carbon double bonds in the vegetable oils (double-bonded carbons bond to fewer hydrogen atoms than single-bonded carbons — i.e., they are unsaturated with respect to hydrogen), converting them into saturated fats, which are generally solids at room temperature.  Another use for hydrogen is in rocket fuels:  the Saturn V rockets that launched the Apollo lunar missions used 209,000 gallons of kerosene and 334,500 gallons of liquid oxygen in its first stage (S-IC), 260,000 gallons of liquid hydrogen and 83,000 gallons of liquid oxygen in its second stage (S-II), and 69,500 gallons of liquid hydrogen and 20,150 gallons of liquid oxygen in its third (S-IVB) stage; the Space Shuttle main engines use 385,000 gallons of liquid hydrogen and 143,000 gallons of liquid oxygen.

Hydrogen is lighter than air, and was used in balloons and dirigibles (also known as airships or zeppelins).  Dirigibles were used in city-to-city air travel in the early 1900s, and in trans-Atlantic crossings in the 1920s and 1930s.  (During World War I, German zeppelins were used in bombing runs over England, since they could fly higher than the British planes.)  On May 6, 1937, the German dirigible Hindenburg caught fire as it came in for a landing at Lakehurst Naval Air Station in New Jersey; 35 people out of the 97 aboard and one person on the ground were killed.  The exact cause of the fire is still the subject of speculation, but the disaster signaled the beginning of the end for airship travel.  Modern "blimps" use helium to provide lift, which avoids the problem of hydrogen's flammability.

Molecules which contain hydrogen bonded to nitrogen, oxygen, or fluorine can attract one another through the formation of hydrogen bonds.  Hydrogen bonds are a particularly strong form of dipole-dipole forces, which arise because of the unequal sharing of electrons in some covalent bonds.  If one atom in a covalent bond is more electronegative than the other, it "pulls" harder on the electrons that the two atoms share, giving the more electronegative atom a partial negative charge, and the less electronegative atom a partial positive charge.  The partially negative atom on one molecule attracts the partially positive atom on a neighboring molecule, causing the two molecules to be more attracted to each other than two nonpolar molecules (which have no electronegativity differences between their bonded atoms) would be.  Molecules that interact by these dipole-dipole forces tend to have higher boiling points than nonpolar molecules, because higher temperatures are necessary to overcome the attractive forces between the molecules and separate the molecules into the gas phase.  In the case of O—H, N—H, and F—H bonds, the electronegativity differences are particularly large because fluorine, oxygen, and nitrogen are the most strongly electronegative elements.  The attractive forces between molecules containing these bonds are particularly strong, and are given the name hydrogen bonds.  Hydrogen bonds are not as strong as covalent bonds, but they greatly influence the physical properties of many substances.  In particular, hydrogen bonds are responsible for the fact that water is a liquid at temperatures at which molecules of similar molecular mass are gases.  (For instance, hydrogen sulfide, H2S, which weighs 34.08 g/mol, boils at -60.28C, while water, weighing in at a measly 18.02 g/mol, boils at 100C.)  Ice floats on liquid water because the hydrogen bonds hold the molecules into a more open, hexagonal array, causing the solid form to be less dense than the liquid form.  In living systems, hydrogen bonding plays a crucial role in many biochemical process, from the coiling of proteins into complex three-dimensional forms to the structure of the DNA double helix, in which the two strands of DNA are held together by the hydrogen bonding between their nucleic acids components.

Hydrogen is also important in a form of spectroscopy called Nuclear Magnetic Resonance (NMR).  In this technique, a sample is placed in a powerful magnetic field (usually produced by a superconducting magnet — see the section on Helium), which causes the hydrogen atoms in the sample to resonate between two different magnetic energy levels; pulsing the sample with a burst of radiofrequency radiation (typically between 200 to 500 MHz) causes the hydrogen atoms to absorb some of this radiation, producing a readout called an "NMR spectrum" which can be used to deduce a great deal of structural information about organic molecules.  Since almost all organic molecules contain hydrogen atoms, this technique is widely used by organic chemists to probe molecular structure; it can also be used to determine a great deal of information about extremely complex molecules such as proteins and DNA.  The technique is nondestructive, and only requires small amounts of sample.  NMR spectroscopy can also be performed with the carbon-13 isotope, and several other isotopes of other elements.  This technology is also used in an important medical imaging technique called Magnetic Resonance Imaging (MRI); the water molecules in different environments in the body respond to very slightly different magnetic field strengths, allowing images of tissues and organs to be obtained.  This technique can be used in diagnosing cancers and creating images of tumors and other diseased tissues.  MRI is also used to study how the brain works by looking at what areas of the brain "light up" under different stimuli.  (The term "nuclear" is avoided in the medical application because of its unpleasant associations, even though the only radiation involved is similar to that of an FM radio transmitter).


Lithium (Li, Z=3).

Lithium is a soft, silvery metal, with a very low density, which reacts vigorously with water, and quickly tarnishes in air.  The name of the element is derived from the Greek word for stone, lithos.  It is found in the Earth's crust at a concentration of 20 ppm, making it the 31st most abundant element.  It is found in the ores spodumene [lithium aluminum inosilicate, LiAl(SiO3)2], petalite [lithium aluminium tectosilicate, LiAlSi4O10], lepidolite [KLi2Al(Al,Si)3O10(F,OH)2] and amblygonite [(Li,Na)AlPO4(F,OH)].

Lithium also presents some exceptions to the "typical" Group 1A behaviors.  The lithium ion has a very high charge density because of its small size; thus, many lithium salts have significant covalent-bonding character, instead of being purely ionic.  These salts dissociate less easily in water than the salts of sodium and potassium, and are therefore less soluble in water.  In addition, lithium can form bonds to carbon which have high covalent character (the organolithium compounds).  Lithium was one of the three elements produced in the Big Bang, although it was produced only in trace amounts. 

Aluminum and magnesium alloys of lithium are strong and lightweight; aluminum-lithium alloys are used in aircraft construction, trains, and bicycles.  Lithium-based batteries have very long lifetimes (particular important in implantable devices such as pacemakers and defibrillators), and are very lightweight; they are frequently used in portable electronic devices and computers.

Lithium salts (such as lithium carbonate, Li2CO3) are used in the treatment of bipolar disorder and some types of depression, and are also used to augment the actions of other antidepressants.  Lithium deuteride (LiD, see entry on Hydrogen above) is used in hydrogen bombs; neutrons produced by a fission-powered explosive are absorbed by the lithium atoms, transforming them into tritium; the fusion of tritium and deuterium to form helium releases tremendous amounts of energy.  Lithium hydroxide (LiOH) is used in confined spaces to remove carbon dioxide from the air (the carbon dioxide is captured in the form of lithium carbonate); this is particularly important in submarines and spacecraft.  [The improvised device for adapting the LiOH canister from the Command Module section of the Apollo spacecraft to fit in the Lunar Module section provided some tension during the ill-fated Apollo 13 mission, and was featured in the movie Apollo 13 (1995).]


Sodium (Na, Z=11).

Sodium is a soft, silvery metal that reacts very vigorously with water, and tarnishes easily in air.  It is the fourth most abundant element in the Earth's crust, which consists of 2.6% sodium by weight; seawater is about 1.5% sodium.  The name is derived from the English word soda, a term found in many compounds of sodium, such as washing soda (sodium carbonate or soda ash), sodium bicarbonate (baking soda), and sodium hydroxide (caustic soda).  The symbol "Na" is derived from the Latin name for the element, natrium.  It is found in the minerals halite [rock salt, or sodium chloride, NaCl] and trona [sodium carbonate bicarbonate, Na3(CO3)(HCO3)], and can be extracted from seawater.  Of the salt that is obtained from these sources, 60% is converted to sodium hydroxide, chlorine, or sodium carbonate; another 20% is used in the food industry as a preservative and flavoring agent, and another 20% is used for other applications, such as de-icing roads.  Metallic sodium is usually stored in mineral oil or some other hydrocarbon, because it will react with the moisture in the air to form sodium hydroxide.

A common laboratory demonstration illustrates the reactivity of sodium.  A small piece of sodium placed in a dish of water skates around on the surface of the water, hissing violently, and slowly disappears.  The sodium reacts with water in a single-displacement reaction, producing sodium hydroxide and hydrogen gas:

2Na(s)  +  H2O(l)    NaOH(aq)  +  H2(g)

The sodium hydroxide is soluble in water, and dissolves.  This demonstration can become very dangerous if too large a piece of sodium is used, however, since enough heat can be generated to ignite the hydrogen gas.

Sodium also reacts vigorously with chlorine gas, producing sodium chloride:

2Na(s)  +  Cl2(g)    2NaCl(s)

This reaction releases a great deal of heat energy, and is usually done in a beaker lined with sand to prevent the heat from cracking the glass.  (See here for a demonstration.)

Energetically excited sodium atoms glow with a yellow light (the strongest emissions are the "sodium D-lines" at 589.0 and 589.5 nanometers), and are prominent in the light from many stars (including the Sun).  Sodium is also used in sodium-vapor street lamps.

In the body, sodium ions regulate osmotic pressure and blood pressure, and sodium and potassium ions together play a major role in the transmission of nerve impulses.

One of the most important compounds of sodium is sodium chloride, NaCl, also known as table salt.  Commercially prepared sodium chloride is either mined in the form of halite, from deposits formed by ancient, dried-out sea beds, or by the evaporation of water from sea water.  Sodium chloride is subjected to electrolysis in an apparatus called a Downs cell, which produces sodium metal and chlorine gas; the construction of the cell is designed to keep the sodium and chlorine separate from each other as they are produced.  Sodium carbonate, Na2CO3, also known as soda or soda ash, has been used for centuries in washing clothes (it helps to remove highly charged metal cations, such as calcium and magnesium, from hard water) and in the manufacture of glass, paper, and detergents.  Sodium hydroxide, NaOH, also known as caustic soda or lye, is a strong base; it is used in drain cleaners, and in the manufacture of detergents (sodium hydroxide breaks down triglycerides — fats and oils such as lard, shortening, olive oil, vegetable oils, etc. — to produce carboxylate salts that form effective soaps).  Sodium bicarbonate, NaHCO3, also known as sodium hydrogen carbonate, is the main ingredient in baking soda, and is used as a leavening agent in the making of bread and other baked goods.


Potassium (K, Z=19).

Potassium is a soft, silvery metal that reacts extremely vigorously with water, and tarnishes rapidly in air.  Its name is derived from the English word "potash," for potassium carbonate, a compound found in high concentrations in wood ashes.  The symbol "K" is derived from the Latin name for the element, kalium.  Potassium is the eighth most abundant element in the Earth's crust (2.1%).  The main ores in which potassium is found are sylvite [potassium chloride, KCl], carnallite [KMgCl36H2O], and alunite [KAl3(SO4)2(OH)6].

Potassium is essential for plant growth, and is heavily used in fertilizers.  In the body, potassium plays a vital role in the contraction of muscle tissue; the movement of sodium and potassium ions in nerve cells plays a major role in the transmission of nerve impulses.  When heated, potassium salts glow with a purple color, and are used in fireworks.  Like sodium, metallic potassium is usually stored under mineral oil or some other hydrocarbon; it can also react with oxygen in dry air to produce potassium superoxide, KO2 (see below).

Potassium undergoes a a reaction with water similar to that of sodium; the products of the reaction are potassium hydroxide and hydrogen gas.  This reaction releases a great deal of heat energy, often igniting the hydrogen gas that is produced.

Potassium-40, which accounts for 0.0117% of the world's potassium, is radioactive, with a half-life of 1.25 billion years.  It undergoes electron capture to produce argon-40; a comparison of the ratio of potassium-40 to argon-40 in rocks can be used to determine the age of the rock (potassium-argon dating).  Trace amounts of potassium-40 are found in all sources of potassium; in a typical human, about 170,000 atoms of potassium-40 decay every second.  The energy released by the decay of potassium-40 is partially responsible for the interior heat of the Earth, along with the decays of thorium and uranium.

There are a number of widely-used compounds of potassium.  Potassium chloride, KCl, is used in salt substitutes (mixed with sodium chloride to improve its flavor), and in fertilizers; massive amounts of potassium chloride are used in lethal injections to cause rapid death by cardiac arrest.  Potassium carbonate, K2CO3, also known as potash, is used in the manufacture of glass.  Potassium hydroxide, KOH, also known as caustic potash,  is used in making soaps and detergents.  Potassium nitrate, KNO3, also known as saltpeter, is a powerful oxidizer, and is one of the ingredients of gunpowder.  Potassium chlorate, KClO3, is a very powerful oxidizer, and is used in match heads and fireworks.  Potassium superoxide, KO2, reacts with carbon dioxide to produce potassium carbonate and oxygen gas; it is used in rebreathers and respiration equipment to generate oxygen, and is also used in mines, submarines, and spacecraft.


Rubidium (Rb, Z=37).

Rubidium is a soft, white metal; it is similar to sodium and potassium in its reaction with water, but the reaction is even more violently exothermic.  Its name is derived from the Latin word for deep red (ruby), rubidius.  It is found in the Earth's crust at a concentration of 90 ppm, making it the 22nd most abundant element.  It is not found in any unique minerals, but is present in trace amounts in lepidolite, pollucite, carnallite, zinnwaldite, and leucite.

Rubidium melts at 39C (102F), so in Texas (where I am writing this) it may be a liquid instead of a solid if the air conditioning isn't working that day.  Metallic rubidium spontaneously combusts in air.  In flame tests, rubidium salts produce a reddish-violet color, and are sometimes used in fireworks.  Rubidium is used in the manufacture of vacuum tubes and cathode ray tubes (CRTs), and is used in some atomic clocks.

In 1995, a vapor consisting of 2000 rubidium-87 atoms was cooled to 170 nanokelvins (17010-9 degrees above absolute zero), producing the first Bose-Einstein condensate, a bizarre state of matter in which all of the atoms occupy the same quantum state, effectively acting as a single superatom (Nobel Prize in Physics, 2001).


Cesium (Cs, Z=55).

Cesium (also spelled as "caesium") is silvery-gold colored metal, which melts at 28C (82F); a sample of cesium will melt in your hand (not that I'd recommend doing this!).  Cesium undergoes the same reaction in water as lithium, sodium, and potassium, but even more violently; because cesium is a very large atom, the outermost electron is lost very easily, and the process is extremely exothermic.  The name is derived from the Latin word caesius, which means "sky blue," because salts of cesium produce a blue color when heated.  Cesium is found in the Earth's crust at a concentration of 3 ppm, making it the 46th most abundant element.  The main ore of cesium is pollucite [CsAlSi2O6]; the refining of pure cesium is made even more difficult by the presence of trace amounts of rubidium in the ore, which is chemically very similar to cesium and thus difficult to separate. 

Because cesium is so reactive, it is used as a "getter" to remove all traces of other gases from vacuum chambers, cathode ray tubes, and vacuum tubes.  Some cesium salts give off light when exposed to X-rays and gamma rays; they are also used in photoelectric cells.

Cesium is used in atomic clocks.  In the SI system, a second is defined as 9,192,631,770 cycles of the radiation corresponding to the energy difference between the ground state and one of the excited states of the cesium-133 atom.

Radioactive cesium-137 is produced in the testing of nuclear weapons, and in nuclear power plants; the explosion at the Chernobyl power plant in 1986 released large amounts of cesium-137, which contaminated a great deal of Western Europe.  Cesium-137 has a half-life of 30 years, and undergoes beta-decay to produce barium-137m, a metastable isotope of barium with a half-life of 2.6 minutes, which emits gamma rays to produce stable ground-state barium-137.

Since cesium ions are so heavy, research on the use of cesium in ion propulsion drives aboard spacecraft and satellites is being conducted.


Francium (Fr, Z=87).

Francium is an extremely rare, radioactive metal.  Its is named for France, the country in which it was first isolated.  It is found in the Earth's crust only in trace amounts, and is one of the least abundant elements on the Earth.  Traces of it are found in uranium ores, where it is produced in the decay series of uranium-235; there is probably only about 20 to 30 grams of naturally-occurring francium in the entire Earth.

All of the isotopes of francium are radioactive, and most have half-lives of less than five minutes; the longest-lived isotope (francium-223) has a half-life of 21.8 minutes.

The possible existence of francium was predicted by Mendeleev from a gap in his periodic table, but the element wasn't discovered until 1939, by Marguerite Perey, an assistant to Marie Curie at the Radium Institute in Paris.




John Emsley, The Elements, 3rd edition.  Oxford:  Clarendon Press, 1998.

John Emsley, Nature's Building Blocks:  An A-Z Guide to the Elements.  Oxford:  Oxford University Press, 2001.

David L. Heiserman, Exploring Chemical Elements and their Compounds.  New York:  TAB Books, 1992.